Preparing and Diluting Solutions
Concentration and Absorbance
Introduction
Solutions are in an important part of chemistry. But how are accurate concentrations
of solutions prepared? In this laboratory activity, a copper (II) sulfate solution will first
be prepared, then diluted to prepare several other solutions with different
concentrations.
Concepts
Concentration
Dilution
Absorbance
Molarity
Background
The amount of solute that is dissolved in a given quantity of solvent is called the
concentration of the solution. A dilute solution contains only a small amount of solute
in a given amount of solution, while a concentrated solution contains a large amount of
solute in a given amount of solution. Molarity is the unit most commonly used to
describe the concentration of a solution. The molarity, M, of a solution is defined as
the number of moles of solute in one liter of solution (Equation 1).
Molarity= amount of solute in moles
Volume of solution in liters
Equation 1
Combined with the molar mass of a solute, Equation 1, used to calculate the number
of grams of solute needed to prepare a given volume of a solution with a specific
concentration. For example, consider the preparation of 500.0mL of a 0.80M solution
of sodium chloride:
Grams of NaCl= 500.0mL x
1L x 0.80moles NaCl x 58.5 g NaCl = 23 g NaCl
1000mL
1L NaCl
1mole NaCl
The calculations show that 23g of sodium chloride are required to prepare 500.0mL of
0.80M NaCl solution.
Once the calculations have been done to determine how much solute is needed to
prepare a solution, precise analytical techniques must be followed to ensure accuracy
in making the solution. Part A, in the Procedure section, describes in detail the proper
analytical procedure for preparing a solution.
One very important aspect of analytical technique involves choosing the right type of
glassware. Volumetric glassware is glassware that has been calibrated (and marked)
to hold a specific volume. The most common form of volumetric glassware used for
preparing solutions is the volumetric flask (figure 1), which has a long, narrow neck
with a single, hairline marking on it. For a 100-mL volumetric flask, the mark on the
neck indicated that when filled to the mark, the flask will contain precisely 100.0 mL at
room temperature.
Diluting Solutions
Experiments often require a solution that is more dilute than what is on hand in the
stockroom. In this case, a more concentrated stock solution must be diluted to obtain
the desired concentration. To carry out a dilution, the following equation is used:
Molarity concentrated soln X volume concentrated soln= Molarity dilute soln X volume dilute soln
In this equation, Molarityconcentrated soln is the concentration of the stock solution,
volumeconcentrated soln is the volume of the stock solution required to prepare the dilute
solution, Molaritydilute soln is the concentration of the desired dilute solution, and
volumedilute soln is the volume of the dilute solution needed. The dilution equation is
commonly written as shown in Equation 2. The subscripts 1 and 2 refer to the
concentrated solution and the dilute solution, respectively.
M1V1=M2V2
Equation 2
For example, assume that the 0.80M sodium chloride solution prepared in the example
above is in the stockroom, but for another experiment, 100mL of a 0.20M sodium
chloride solution is needed. In performing a dilution calculation, M1, M2, and V2 are
generally known and Equation 2 is rearranged to solve for the unknown V1.
Substituting the known values for this example into Equation 2 allows us to solve for
the volume of the concentrated solution required to prepare the dilute solution.
V1=M2V2
M1
V1=0.20M * 100mL = V1=25mL
0.80M
Proper analytical technique for preparing the diluted solution requires that the initial
and final volumes (V1 and V2) must be accurately measured using a graduated cylinder
or, preferably, a pipet and a volumetric flask.
Concentration and Absorbance
Molarity and dilution calculations show us how to prepare solutions of known
concentration. Another important problem chemists encounter in the lab is how to
determine the concentration of an unknown solution. If the solution is colored, the
concentration of an unknown solution can be determined by measuring the intensity of
the color. A special sensor or instrument called a colorimeter is used to measure the
absorbance of visible light that gives the solution its color. Generally, the more intense
the color of the solution, the greater the absorbance of light will be. In using
colorimetry, it is important to remember that the color of light transmitted by the
solution (the color we see) is complimentary to the color of light absorbed by the
solution (the color we measure). Since the color of light depends on its wavelength,
the wavelength of light absorbed by a colored substance in solution is complementary
to the wavelength of light transmitted by the substance. Copper (II) sulfate solutions,
for example, are blue. The absorbance of copper (II) sulfate solutions is measured at
635nm, corresponding to red light.
Experiment Overview
The purpose of this experiment is to prepare a series of blue copper (II) sulfate
solutions of known concentration using the molarity and dilution equations. The
relationship between the concentration of a solution and its absorbance will be
investigated. The accuracy of the solution preparation and dilution procedures will
then be determined.
Pre-Lab Questions
1. Calculate the number of grams of copper (II) sulfate pentahydrate, CuSO4 • 5H2O,
required to prepare 100.0mL of a 0.150M copper (II) sulfate solution.
2. Calculate the number of milliliters of 0.150M copper (II) sulfate solution that must
be diluted to prepare 10.0mL of a 0.0750M copper (II) sulfate solution.
3. Calculate the number of milliliters of 0.150M copper (II) sulfate solution that must
be diluted to prepare 10.0mL of a 0.0230M copper (II) sulfate solution.
Materials
Copper (II) sulfate pentahydrate,
CuSO4 • 5H2O, 3g
Balance, centigram precision
Beral-type pipet
Colorimeter or spectrophotometer
Cuvets, 5
Graduated cylinder, 10-mL
Spatula
Test tube rack
Test tubes, 5
Tissues of lens paper
Volumetric flask, 100-mL
Wash bottle
Wax pencil or labeling tape
Weighing dish
Safety Precautions
Copper (II) sulfate is moderately toxic by ingestion and inhalation and is a skin and
respiratory irritant. Avoid contact with eyes and skin. Wear chemical splash goggles,
chemical resistant gloves, and a chemical- resistant apron. Wash hands thoroughly
with soap and water before leaving the laboratory.
Procedure:
Part A. Preparing the Stock Solution
1. Review the calculations from Pre-Lab Question #1 for the number of grams of
copper (II) sulfate pentahydrate, CuSO4 • 5H2O, required to prepare 100.0 mL of a
0.150 M solution. Check them with your instructor. Once your calculations have
been approved, weigh out the required amount of copper sulfate on a balance in a
clean, dry weighing dish.
2. Transfer the solid to a 100-mL volumetric flask.
3. Use a wash bottle filled with distilled or deionized water to rinse any remaining solid
from the weighing dish into the flask.
4. Add more distilled or deionized water to the volumetric flask until the liquid level is
almost to the 100-mL mark. Fill to the mark with a pipet or wash bottle drop-bydrop so that no water splashes up on the sides of the flask. Fill until the bottom of
the meniscus is EXACTLY at the 100.0-mL mark.
5. Carefully add a stir bar to the solution in the flask. Place the flask in the center of
the stirrer/hotplate to mix the solution.
6. Stir the solution to give a homogenous solution.
Part B. Preparing Diluted Solutions
7. Place five clean and dry test tubes in a test tube rack and label them #1-5. Label
one pipet “CuSO4” and use it to transfer the stock solution only.
8. Using a 10mL graduated cylinder, measure and pour 10mL of the 0.150M stock
solution into test tube #1. Record the necessary data for this solution in the data
table.
9. Using a clean Beral-type pipet, fill the 10-mL graduated cylinder exactly to the 3.80mL mark with the stock solution. Try not to get any drops of solution on the sides
of the cylinder. Make sure that the bottom of the meniscus sits exactly at the 3.80mL mark.
10. Carefully fill the graduated cylinder to the 10.0-mL mark with distilled or deionized
water. Do not overfill!
11. Mix the solution in the graduated cylinder by repeatedly filling and emptying the
test tube with the solution three times. The agitation caused by filling and emptying
the test tube will mix the solution.
12. Transfer the mixed solution to test tube #2 and record the necessary data for this
solution in the data table.
13. Rinse the graduated cylinder with water and dry it with a paper towel.
14. Repeat steps 9-11 using 2.40mL stock solution.
15. Transfer the mixed solution to test tube #3 and record the necessary data for this
solution in the data table.
16. Rinse the graduated cylinder with water and dry it with a paper towel.
17. Before proceeding, review your calculations for preparation of 10.0mL of a 0.075M
and a 0.023M cupric sulfate solution (Pre-Lab Questions #2 and #3). Check with
your instructor before proceeding.
18. Using your calculations and the analytical technique described in steps 9-13,
prepare 10.0mL of a 0.075M cupric sulfate solution by diluting the stock solution.
Transfer this solution to test tube #4 and record the necessary data in the data
table.
19. Using your calculations and the analytical technique described in steps 9-13,
prepare 10.0mL of a 0.023M cupric sulfate solution by diluting the stock solution.
Transfer this solution to test tube #5 and record the necessary data in the data
table.
20. Compare the color of the stock solution and each of the dilutions in test tubes #15. Rank them in terms of color from deepest blue to lightest blue. Record these
observations in the data table.
Part C. Colorimetry Measurements
21. Rinse five cuvets with about 1mL of the solutions from Part B, and then fill the
cuvets with each solution. Arrange the cuvets in order (#1-5) on a labeled sheet of
paper to keep track of the solutions. Do not write on the cuvets.
22. Handle the cuvets by their ribbed sides or their top to avoid-getting fingerprints on
the surface. Wipe the cuvets with lint-free paper or lens paper.
23. Connect the interface system to the computer or calculator and plug the
colorimeter sensor into the interface.
24. Select Setup and Sensors from the main screen and choose “Colorimeter”. Note:
Many newer sensors have an automatic calibration feature that automatically
calibrates the colorimeter before use. If the sensor has the auto calibration
feature, proceed mission (0 absorbance) with a “blank” cuvet containing only
distilled water.
25. Select Calibrate and Perform Now from the Experiment menu on the main screen.
26. Fill a cuvet about ¾ full will distilled water. Wipe the cuvet with a lint-free tissue,
then place the cuvet in the colorimeter compartment.
27. Set the wavelength knob on the colorimeter to 0%T--the onscreen box should read
zero. Press Keep when the voltage is steady.
28. Turn the wavelength knob on the colorimeter to 635nm (red)--the onscreen box
should read 100. Press Keep when the voltage is steady.
29. Return to the main screen and set up a live readout and data table that will record
absorbance (as a function of time).
30. Select Setup followed by Data Collection. Click on Selected Events to set the
computer for manual sampling.
31. Remove the “blank” cuvet from the colorimeter compartment and replace it with
the cuvet containing solution #1.
32. Press Collect on the main screen to begin absorbance measurements.
33. When the absorbance reading stabilizes, press Keep on the main screen to
automatically record the absorbance measurement. Note: The absorbance
measurement should appear in a data table onscreen. The onscreen data table will
also contain a time reading, which may be ignored.
34. Remove the cuvet from the colorimeter compartment and replace it with the cuvet
containing solution #2.
35. When the absorbance reading stabilizes, press Keep on the main screen to
automatically record the absorbance measurement.
36. Repeat steps 34 and 35 with the other solutions #3-5.
37. Press Stop on the main screen to end the data collection process. If possible,
obtain a printout of the data table.
38. Record the absorbance data for solutions #1-5 in the Data Table.
39. Dispose of the contents of the cuvets and of the remaining test solutions as
directed by your instructor.
40. Follow your instructor’s directions for rinsing and drying the cuvets.
Data Table
Test Tube
Volume of Stock Solution (V1)
1
2
3
4
5
Concentration of Stock
Solution (M1)
Final Volume of Diluted
Solution (V2)
Concentration of Diluted
Solution (M2)
Color Comparison
(Rank Solutions from lightest
blue=1, deepest blue = 5)
Absorbance at 635 nm
Post-Lab Questions
1. Calculate the concentrations of the diluted solutions in test tubes #2 and 3 using
Equation 2 from the Background section.
2. Complete the data table for test tubes #1, 4, and 5. Note: See the Pre-Lab
Questions for the calculated values for solutions #4 and 5.
3. Compare the concentration of each solution to the color ranking. What is the
relationship between the concentration of a solution and its color intensity?
4. Compare the concentration of each solution with its absorbance. What is the
relationship between concentration of a solution and its absorbance?
5. Prepare a graph of absorbance on the y-axis versus the concentration of each
solution on the x-axis.
6. Does it make sense that the relationship between concentration and absorbance
should include the origin (0,0) as a point? Explain your reasoning.
7. Based on your answer to Question #6, draw the “best-fit” straight line through the
data points. How well does this straight line fit the data? Describe the accuracy of
the relationship between concentration and absorbance.
8. The absorbance of a copper (II) sulfate solution of unknown concentration was
measured by colorimetry and found to be 0.250. Use your graph of absorbance
versus concentration to estimate the concentration of this unknown solution.